Chapter 17: Inorganic Chemistry

The s-block

Group 1 (alkali metals) are soft, reactive metals with one outer electron; reactivity increases down the group: . Group 2 (alkaline earth metals) are less reactive. Down Group 2, the thermal stability of carbonates and nitrates increases, the solubility of hydroxides increases, and the solubility of sulfates decreases. These trends are explained by the polarizing power of the cation: smaller, more highly charged cations distort the anion more and so decompose more readily. Flame colours help identify ions (Na yellow; K lilac; Ca brick-red; Sr crimson; Cu blue-green).

The p-block

Group 13: aluminium is amphoteric (its oxide and hydroxide react with both acids and bases) and is extracted by electrolysis of alumina dissolved in molten cryolite (the Hall–Héroult process).

Group 14: carbon (diamond, graphite, fullerenes), silicon (semiconductors, silica, silicates).

Group 15: nitrogen is unreactive because of its strong N≡N bond. Ammonia is made by the Haber process: N₂ + 3H₂ ⇌ 2NH₃ (), using an iron catalyst at about 450 °C and 200 atm, a compromise between rate and yield. Nitric acid is made by the Ostwald process.

Group 16: oxygen, and sulfur with its oxides and acids; sulfuric acid is made by the Contact process (SO₂ + ½O₂ ⇌ SO₃ over V₂O₅).

Group 17 (halogens): oxidizing power decreases down the group (Cl₂ displaces Br⁻ and I⁻); the acid strength of the hydrogen halides increases HF < HCl < HBr < HI because the H–X bond becomes weaker; the silver halides vary in solubility in ammonia.

Group 18: noble gases are largely inert; xenon forms fluorides such as XeF₂ and XeF₄.

The d-block

The transition elements form ions with partly filled subshells. Characteristic properties: variable oxidation states, coloured compounds, catalytic activity, and formation of complex ions. A complex has a central metal ion surrounded by ligands that donate lone pairs (coordinate bonds). The coordination number is the number of donor atoms (commonly 4 or 6). Typical geometries: octahedral (6), tetrahedral and square planar (4). Examples: (blue), (deep blue), . Crystal field theory explains colour: ligands split the orbitals into sets of different energy, and the ion absorbs light of energy equal to the gap, so its colour is the complement of the colour absorbed.

Redox titration example. In acid, MnO₄⁻ oxidizes Fe²⁺: MnO₄⁻ + 5Fe²⁺ + 8H⁺ → Mn²⁺ + 5Fe³⁺ + 4H₂O. If 20.0 cm³ of 0.0200 mol dm⁻³ KMnO₄ is needed, moles of MnO₄⁻ , so moles of Fe²⁺ .

Qualitative analysis

Always follow the exact tests and safety rules given in your laboratory manual and the syllabus.

Industrial and environmental chemistry

Large-scale processes balance rate, yield, cost, safety, and environmental impact: Haber (ammonia), Contact (sulfuric acid), blast-furnace extraction of iron, electrolytic extraction of aluminium, and the chlor-alkali process. Environmental topics include the greenhouse effect (CO₂ and CH₄ absorbing infrared radiation), acid rain (SO₂ and NOₓ), ozone depletion (CFCs), water treatment, and green chemistry principles such as waste prevention and atom economy.

Common mistakes

Memorizing trends without the explanation (polarizing power, bond strength, lattice enthalpy).

Confusing ligands with counter-ions when naming and counting charge.

Using wrong conditions or catalysts in industrial processes.

Practice questions

Explain why barium carbonate is more thermally stable than magnesium carbonate.

Write the oxidation number of Mn in KMnO₄ and in MnO₂. [+7; +4]

What is the coordination number and oxidation state of Fe in ? [6; +3]